Download ATOMIC STRUCTURE

ATOMIC STRUCTURE
[MH5; Chapter 6]
• An atom consists of:
• In a neutral atom:
• As we look at the arrangements of electrons in atoms, we are paying
special attention to:
1) the relative energies of the electrons
2) the spatial locations of the electrons
• The electron configuration (the actual arrangement of the electrons)
of an element can be determined from its position on the Periodic
Table.
• Recall that we have already used these positions can also be used to
predict sizes of atoms and their tendencies to gain or lose electrons.
-101-
Light, Photon Energies and Atomic Spectra
[MH5; 6.1]
The Wave Nature of Light
• Light travels through space as a wave; the wavelength, λ , is the
distance between successive crests or troughs.
• The number of cycles that pass a given point in unit time is the
frequency, ν , of the wave.
• A frequency of one Hertz, Hz, is one cycle per second.
• The speed of light, c, the speed at which a light wave moves
through space is a constant,
2.998 × 108 m sG1.
• The eye sees only a small fraction of the electromagnetic spectrum,
the Visible region, which has wavelengths ranging from 400-700 nm.
• Ultra violet (UV) and Infra red (IR) are two of several other regions.
-102-
The Particle Nature of Light
• Work by Max Planck and Albert Einstein during the years 1900-1910
showed that light is generated as a stream of particles, called
photons, whose energy is given by the Einstein equation:
• The quantity h is called Planck’s constant, with a value of
h = 6.626 × 10G34 J s
• Note that the energy of the photon is inversely proportional to
wavelength, and directly proportional to frequency.
-103-
Atomic Spectra
• White (or visible) light (from the sun) can be broken down into its
colour components; this is called a spectrum.
• This spectrum is continuous and contains essentially all wavelengths
between 400 and 700 nm.
• If an element is vaporized and then excited (given more energy);
photons of light are emitted as the excited electrons return to
their pre - excited energy level....this is called an atomic spectra.
• This is not a continuous spectra; the emitted light creates a line
spectra in which each line corresponds to a specific wavelength of
light.
• Atomic spectra are characteristic to the element in question; can be
used for identification.
• As photons have only certain wavelengths, it follows that they can
only have certain energies.
• As photons are produced when an electron moves from one energy
level to another; it follows that the energy levels for an electron in
an atom are limited to specific values.
• So we say that the energy of an atom is quantized......it exists in
certain fixed quantities.
The Hydrogen Atom
• Hydrogen atoms are produced when an electric discharge is struck
through gaseous hydrogen [ H2 (g) ! 2 H* (g, at)]
•
•
These gaseous hydrogen atoms emit radiation (light) at wavelengths
which may be grouped into series.
The first of these series to be discovered was the Balmer Series,
which lies partly in the visible region.......
Red, 656.28; Green, 486.13; Blue violet, 434.05 and Violet, 410.18.
-104-
•
So we have seemingly contradictory results...........
1) the continuous spectrum obtained from white light
2) the line spectrum obtained from atoms
The Bohr Model of the Hydrogen Atom
• This model of the Hydrogen atom is a theoretical explanation of
atomic spectra; based on classical mechanics.
• Bohr assumed a central proton, about which an electron moves in a
circular orbit.
• The electrostatic attractive force between the positively charged
proton and the negatively charged electron is balanced by the
centrifugal force due to the circular motion.
• Energy would then be absorbed/radiated as the electron changed
the radius of its orbit.
-105-
•
•
•
Such a model predicts a continuous range of possible electron
energies, which would give rise to a continuous rather than a line
spectrum.
Bohr boldly assumed that only certain definite energies of the
electron were allowed.
He had no theoretical basis for this assumption, merely the need to
explain the observed experimental result. He deduced:
En = !RH / n2
where RH = Rydberg constant = 2.180 × 10!18J, and
n = integer = 1, 2, 3...
•
This integer, n, becomes known as the principal quantum number, or
energy level of the electron.
•
Bohr’s model established three general points......
1) Zero energy was designated as the point at which electron and
proton are completely separated, i.e. with the electron at infinity.
As the electron approaches the proton, it loses energy.
All allowed electron energy states therefore have negative energies.
2) The hydrogen atom is normally in the ground state, and the electron
has principle quantum number (or energy level) n = 1.
It can absorb energy, and move to an excited state; n > 1.
The first excited state is n = 2; the second excited state is n = 3.....
Transitions are also possible between excited states, eg. from n = 2
to n = 3.
-106-
3) When the electron gives off energy as a photon of light, it falls to a
lower energy state. (eg. from n = 2 to n = 1 or n = 3 to n = 2...)
The energy of the photon evolved (and hence the wavelength/colour)
is equal to the difference in energy between the two states:
∆E = Ehigh ! Elow = hν
Since En = ! RH / n2 the energy of the both states can be calculated,
and thence the wavelength of the photon emitted.
The model worked brilliantly for the hydrogen atom, accounted for
the line spectrum observed, and Niels Bohr received the Nobel Prize
for physics in 1922.
The Quantum Mechanical Model [MH5, p. 136]
• Bohr’s model, although highly successful in explaining the line
spectrum of the hydrogen atom, did not even work for He, much less
more complicated atoms.
• It was soon realized that there was some fundamental problem with
the Bohr model, and the fixed orbit of an electron idea was
abandoned.
•
In 1924 Louis de Broglie suggested that perhaps electrons and light
could behave in a similar manner.......
-107-
•
This new idea (the dual wave - particle nature of the electron) was
experimentally confirmed and led to a whole new discipline commonly
known as quantum mechanics.
•
The Quantum Mechanical Model differs from the Bohr model in
several ways; the two most important being:
1) The kinetic energy of an electron is inversely related to the
volume of the region to which it is confined.
2) It is impossible to specify the precise position of an electron in
an atom at a given instant.
•
•
The region where we expect to find the electron is sometimes called
the electron cloud.
Electron density decreases as the distance from the nucleus
increases.
-108-
•
•
In 1926, Erwin Schrödinger, an Austrian physicist working in Zurich,
wrote a differential equation to express the wave properties of an
electron in an atom.
From this equation, one could calculate the amplitude ψ (wave
function) of an electron wave at various points in space.
The Schrödinger Equation
•
•
•
•
•
•
What is important is that, for a H atom, ψ2 is directly proportional
to the probability of finding the electron at a particular point.
The 3 - dimensional Schrodinger equation (Ψ) can be solved exactly
for the hydrogen atom, and approximately for atoms with 2 or more
electrons.
There are many solutions for Ψ, each associated with a set of
numbers called quantum numbers.
A wave function corresponding to a particular set of three quantum
numbers (n, R, mR; to be defined shortly) is associated with an
electron occupying an atomic orbital.
From these numbers we can deduce the energy of that orbital, its
shape and its orientation in space.
A fourth quantum number, ms, completes the description of a
specific electron.
– 109 –
Quantum Numbers and Electron Configurations [MH5;6.3]
Principal Quantum Number, symbol n
• Derived from the Bohr atom quantum number n, this defines a
principal energy level for the electron in the hydrogen atom.
• n may have any positive integral value > 0, ie) n = 1, 2, 3, 4...
• As n increases, the energy of the electron and its distance from the
nucleus increase.
Orbital Quantum Number, symbol R (Sublevels s, p, d, f)
• Each principal energy level includes one or more sublevels denoted
by the orbital quantum number, R , which determines the general
shape of the electron cloud.
•
R may have integral values from 0 up to (n ! 1); so......
If n = 1, only R = 0 is possible.
If n = 2, R = 0 or R = 1
•
The number of sublevels in a given energy level (up to a point) is
equal to the principal quantum number, n.
•
We have ‘names’ for electrons with different values of R :
R = 0,
R = 2,
‘s electrons’
R = 1,
R = 3,
‘d electrons’
‘p electrons’
‘f electrons’
Sublevel Designations for Principal Quantum Numbers 1 to 4
n
1
2
3
4
R
Sublevel
•
Within a specific energy level (n) , the sublevels increase in energy
in the order:
ns < np < nd < nf
– 110 –
Magnetic Quantum Number, symbol mR
• Each sublevel contains one or more orbitals, which differ in the
direction, or orientation in space taken by the electron cloud.
•
The number mR defines that direction and is related to R.
•
mR may have integral values from !R through zero to +R.
•
If R = 0 (s electron), mR = 0; as there is just one s orbital, there
can only be one possible orientation, or orbital.
•
If R = 1 (p electron), mR = -1, 0 or +1; there are three possible
orientations in a p sublevel, or 3 p orbitals.
Similarly for :
•
R= 2
R= 3
•
All orbitals within a given sublevel have the same energy.
Spin Quantum Number, symbol ms
• This quantum number is associated with spin, possessed by all
electrons (and other fundamental particles).
• Regardless of the values of the other quantum numbers, ms may be
only +½ or -½ .
• These values correspond to a clockwise or counterclockwise spin, but
we don’t have any idea which is which!
• Electrons with the same value of ms (either both + ½ or - ½) are said
to have parallel spins; if the values are different (one at + ½ and one
at - ½) , the electrons are said to have opposite spins.
• Spin is usually represented by arrows pointing up or down.......
  (parallel)
  (opposite)
– 111 –
•
•
•
Recall that when the Schrödinger equation may also be solved for
many-electron atoms, there are many solutions for Ψ, each
associated with a set of quantum numbers.
Each set of four quantum numbers describes a different electron in
the atom.
In 1925, Wolfgang Pauli, a colleague of Bohr, proposed the Exclusion
Principle:
In one atom, no two electrons may have the same four quantum
numbers.
•
The result of this is that only two electrons (with opposite spins)
can fit into any orbital......
Values for Quantum Numbers through n = 4
n
R
mR
ms
1
2
3
4
•
We can use this chart to determine how many electrons we can fit in
each energy level.....this works out to 2n2
– 112 –
Shape of s and p orbitals
•
•
•
•
[MH5; 6.4, Figure 6.7]
s orbitals are spherical; as n increases, the radius of the orbital
becomes larger.
These spheres are give a specific probability of finding the electron.
So an electron in a 2s orbital is likely to be found further away from
the nucleus than an electron in a 1s orbital.
In p orbitals, the R = 1 sublevel, have electron probability density
concentrated along the x, y and z axes:
– 113 –
Electron Configurations in Atoms [MH5; 6.5]
• Using the chart on page 264, we could assign quantum numbers to
each electron in an atom.
• Now we will assign electrons to specific energy levels, sublevels and
orbitals.
• We do this using an Electron Configuration, which shows the number
of electrons (indicated by a superscript) in each sublevel.
• The sum of the number of electrons in each sublevel must be equal
to the total number of electrons in the atom....
• When determining an electron configuration, always begin with the
electrons closest to the nucleus; those at n = 1.
EXAMPLE 1:
H:
He:
Be:
EXAMPLE 2:
B:
N:
CR:
– 114 –
•
In partially-occupied sub-levels, the arrangement of electrons
follows Hund's Rule:
Electrons occupy orbitals of equal energy in such a way that the
number of unpaired electrons is a maximum, thus minimizing
inter-electronic repulsions.
•
This is shown by the arrangements of electrons in partiallyoccupied p orbitals:
p1

p4  


p2 

p5  


p3 


p6  


•
Similarly, the number of unpaired electrons in a d orbital increases
to a maximum of five at 3d5, then decreases as further electrons
are added........
•
Hund's rule tells us the lowest energy, or ground state
arrangement.
Other arrangements consistent with the Pauli principle are
sometimes seen, but are of higher energy, and referred to as
excited states:
•
EXAMPLES:
– 115 –
•
•
•
•
•
•
•
So far, this doesn’t seem too difficult......In many cases, we simply
slot electrons into the orbitals in increasing order of n value and, for
the same n, in increasing order of orbital quantum number, R.
This approach works for the first 18 elements, up to 18Ar,
configuration 1s22s22p63s23p6.
It is also consistent with MH5; figure 6.8.
But after the 3p orbital has been filled, it looks like the 4s orbital is
of lower energy......so it fills next.
In the next two atoms, K and Ca, the ground-state has electrons in
the 4s orbital, giving [Ar] 4s1 and [Ar] 4s2 respectively.
Note the Noble Gas Configuration......one need only write the
configuration past the preceding Noble Gas.
In the next ten elements, Sc to Zn (the first transition series), the
3d orbital is filled. (After 4s, before 4p)
– 116 –
•
•
There are two significant exceptions in electron configurations in
the 3 d subshell; Cr is one of them.
The six electrons over the [Ar] core are arranged 4s13d5 rather
than 4s23d4, both orbitals being partially filled.
•
•
In this arrangement all six electrons are unpaired.
Energies of all the orbitals are approximately equal, and this
arrangement obeys Hund's rule and minimizes inter-electron
repulsions.
•
Towards the end of the transition series, with Cu, the energy of the
3d orbital has fallen below that of the 4s because of the increased
nuclear charge, so the configuration is 3d104s1.
•
Zn is, as expected, 3d104s2. [Similar effects are seen in the second
transition series.]
Relating Quantum Numbers and Electron Configurations
• Now we will actually write some sets of quantum numbers for
electrons.....
EXAMPLE 1:
H:
He:
– 117 –
EXAMPLE 2:
N:
•
We can also determine whether or not a set of quantum numbers is
correct for a particular electron in an atom....
EXAMPLE:
A)
B)
C)
D)
E)
Which of the following is a correct set of quantum
numbers for a highest energy electron in As ?
n
R
mR
ms
3
3
4
4
4
1
2
1
2
0
0
-1
0
+1
0
+½
+½
-½
-½
+½
– 118 –
The Core plus Valence Shell Model
• The chemically important part of an atom is its valence shell,
defined as only those subshells whose electrons influence chemical
properties; valence and oxidation state in particular.
• It is only the valence shell electrons that participate in the
formation of ions; gaining electrons to produce anions or losing
electrons to produce cations.
• The main-group atoms from Group 13 to 17 can be confusing because
the Periodic Table shows a filled d10 subshell from the core coming
between the s and p subshells that constitute the valence shell.
• The transition atoms can also be confusing because here the ns
subshell and the (n – 1) d subshell from the previous energy level are
both part of the valence shell.
EXAMPLES:
A) S
B)
Br
C)
Mn
D)
Sb
•
Determine the valence shell for:
There are eight main groups; for a main-group atom, the number of
electrons in the valence shell is equal to either the group number
[H and He are special cases] or the group number - 10.
– 119 –
valence shell configuration ns1
Group 1
Li, Na, K, Rb, Cs, Fr
Alkali metals
valence shell configuration ns2
Group 2
Be, Mg, Ca, Sr, Ba, Ra
Alkaline Earths
valence shell configuration ns2 np1
Group 13
B, AR, Ga, In, TR
valence shell configuration ns2 np2
Group 14
C, Si, Ge, Sn, Pb
valence shell configuration ns2 np3
Group 15
N, P, As, Sb, Bi
valence shell configuration ns2 np4
Group 16
O, S, Se, Te, Po
Chalcogens
valence shell configuration ns2 np5
Group 17
F, CR, Br, I, At
Halogens
valence shell configuration ns2 np6
Group 18
He (1s2), Ne, Ar, Kr, Xe, Rn
Noble or Rare gases
•
•
•
•
(He is in Group 18 because of its similarity to other atoms having a
filled valence shell).
Groups 1 and 2 are known as s-block atoms and Groups 13 to 18 are
known as p-block atoms.
Orbitals below the valence shell are referred to as ‘core' electrons.
From the Periodic Table, you can determine the valence-shell
electron configuration of any main-group atom from its position.
• Horizontal Rows of the Table may be denoted in two ways:
a) Period
1, 2, 3 ..... corresponding to n = 1, 2, 3 .. etc.
H!He, period 1.
Li!Ne, period 2
etc.
b) Row is assigned differently with H and He ignored. Thus First!row
elements are the series from Li to Ne, second!row elements from
Na to Ar, etc.
– 120 –
Ionic Species [MH5; 6.7]
• Recall that ions are charged species which result from either the
gain or loss of electrons.
• Electron configurations can also be written for ionic species.
• Anions are negatively charged species which are formed when an
atom gains electrons.
• The electron configuration for an anion is the same as the atom
which contains the identical number of electrons as does the anion...
EXAMPLE:
What is the electron configuration for the S2— ion ? The Br — ion?
Cationic Species
• Recall that cations are positively charged ions formed when an atom
loses electrons......
• When writing electron configurations for cations, one must first
determine the orbital(s) from which the electrons were lost....
• The electrons lost when a cation forms are always lost from the
valence shell.
EXAMPLES:
A) Ca2+
Determine the electron configuration for:
B) Bi3+
C) Bi5+
– 121 –
• Isoelectronic species are atoms or ions with the same number of
electrons arranged in the same way.
• The following examples are ions having a noble-gas structure:
EXAMPLES:
O2! F ! Ne Na+ Mg2+ AR3+:
S2– CR! Ar K+ Ca2+ Sc3+ :
all 1s22s22p6 or [Ne]
all 1s22s22p63s23p6 or [Ar]
•
Since all these species have filled levels and sublevels,
configurations characteristic of the unreactive or ‘noble’ gases, it is
not too surprising that they are ‘stable’ species!
•
When a transition metal forms a cation, the electrons are also lost
from the valence shell, which includes the n s orbital and the n - 1 d
orbital........but remember that the electrons are lost from the s
orbital first !!
So....when electrons are lost from a transition-metal atom, giving a
cation, any electrons remaining over the noble-gas core are always
found in the d orbital(s).
•
EXAMPLES:
V atom
Determine the electron configurations for:
!
Vanadium(II) ion, V2+
•
•
Compare this result with the neutral Sc atom.
Like neutral Sc, V2+ has three electrons over the [Ar] core, but the
greater nuclear charge (more protons in nucleus) lowers the energy
of the 3d below that of the 4s orbital....
•
Iron behaves in a similar manner; the neutral Fe atom, [Ar]3d64s2,
– 122 –
loses electrons to form Fe2+ and Fe3+
Fe2+ ion ,
!
Fe3+ ion, + eG
•
•
Note that V2+ and Sc0, or Fe3+ and V0, are NOT isoelectronic pairs.
They do have the same number of electrons over the [Ar] core
(3 and 5 respectively), but the electrons are arranged differently in
the orbitals.
•
This difference is due to the increased nuclear charge (number of
protons is greater than the number of electrons) in the cations.
Electrons in a cation will take the lowest-energy arrangement
possible.
•
ONE MORE EXAMPLE: Determine the electron configuration of:
Mn:
Mn2+:
Mn7+:
– 123 –
The Periodic Table and Electronic Configurations
• The Periodic Table is divided vertically into four blocks according to
the subshell being filled: [see them color-coded in MH5; Figure 6.9]
s-block in blue, p-block in green, d-block in pink, f-block in brown.
•
Main Group atoms are those in which s or p subshells are being
filled, other subshells being either full or empty.
•
Transition elements, sometimes referred to as d-block atoms, are
those in which d orbitals are being filled.
Since each d orbital sublevel holds ten electrons, these fall into
three series, filling respectively the 3d, 4d, and 5d orbitals.
By convention, the last member of each series (Zn, Cd, and Hg) is
thought of as a transition atom, even though the d orbital is full and
is not disturbed in reactions.
•
•
•
The f-block elements, in which f orbitals are being filled. These
orbitals hold a total of 14 electrons, so we have a series of 14 atoms
known as the lanthanides, following (but not including) 57La,
Lanthanum, in which the 4f orbital is filled.
•
A second series following 89Ac, Actinium, contains 14 atoms known as
the actinides, and the 5f orbital is irregularly filled.
It consists entirely of radioactive atoms.
Only 90Th, thorium, and 92U, uranium, occur in nature, the others
have been synthesized in nuclear reactions.
•
•
– 124 –
•
A scheme to help you remember the order for filling the orbitals......
1s
2s
2p
3s
3p
3d
4s
4p
4d
4f
5s
5p
5d
5f
6s
6p
7s
– 125 –
Periodic Properties of Atoms
[MH5; 6.8]
Atomic Size
• The size of the electron cloud is unlimited in extent, so the ‘size’ of
a free atom cannot be exactly defined.
• A quantity called atomic radius, assumes a spherical atom and can
be defined and measured.
• Atomic radius is taken to be one half the shortest distance of
approach between atoms in an elemental substance; this gives us an
effective size......
• In a covalent bonding situation the size is easier to define; the
covalent radius of an atom is the contribution that atom makes to
the internuclear distance (or bond length).
• So the covalent radius of H is half the H—H bond length in H2.
•
•
Ionic radius is used to describe ion sizes; the sum of the ionic radii
equals the interionic distance.
Loss of electrons (giving a cation) always leads to a reduction in size
from the atomic radius:
Li atom
152 pm
!
Li+ cation
+
90 pm radius.
– 126 –
e!
•
A gain of electrons (giving an anion) always leads to an increase in
size:
F atom
72 pm
+
eG
!
F G anion
133 pm radius
•
Assume that nuclear charge dominates in an isoelectronic series....
•
In general, size decreases across a period from left to right and
increases down a group.
These trends can be explained in terms of the effective nuclear
charge, Zeff experienced by the outermost or valence electrons.
For any electron, Zeff is approximately Z ! S, where Z is the actual
nuclear charge (the atomic number) and S is the number of core
electrons. Thus:
•
•
Na
12Mg
13Al
11
•
Only completed shell [core] electrons shield valence electrons from
the nuclear charge; valence electrons do not.
General Trend for Atomic Size in the Periodic Table
– 127 –
Ionization Energy [MH5; page 153]
• Ionization Energy, or I.E., is defined as the energy required to
remove an electron completely (to infinite separation) from an
isolated (gaseous) atom:
!
M(g)
M+(g)
+
eG
•
An atom with several electrons will have successive I.E.’s
First
+
eG
M(g)
!
M+(g)
Second
M+(g)
!
M2+(g)
+
eG
•
Successive I.E. values always increase in magnitude because of the
increasing ionic charge.
Removal of an electron is always an energetically unfavorable
(endothermic) process, but it varies in a periodic manner........
Generally speaking, I.E. increases from Group 1 through Groups 2,
13...18 as the number of electrons in the valence shell increases.
•
•
•
Low I.E.’s are found when the valence shell contains only one
electron shielded from nuclear attraction of the protons by the
lower!energy core electrons.
EXAMPLE:
Li (1s2 2s1)
I.E:
520 kJmolG1
Na (1s2 2s2 2p6 3s1)
I.E:
496 kJmolG1
•
H is an exception with I.E. 1312 kJ molG1 because no core electrons
are present.
Small anomalies are also found
Be (1s2 2s2)
I.E. 900 kJ molG1
B (1s2 2s2 2p1)
I.E. 801 kJ molG1
•
•
•
This occurs because the electron lost in B, Boron, is the 2p1
electron, which is of higher energy than the 2s2.
A higher energy electron is more easily lost.
– 128 –
•
As the next 5 electrons are added to the 2p orbitals, (B through
Ne) the I.E. increases as the nuclear charge increases. But:
N: 2p
O: 2p
•
The electron lost from the O atom is pushed out by the extra
repulsion from the second electron in the same 2px orbital and thus
O has a lower I.E. than N.
•
For elements with similar valence shells, I.E. falls on going down the
group:
Li
520
Na
496
K
419
Rb
403
Cs
377
– 129 –
General Ionization Energy trends in the Periodic Table:
Electronegativity [ MH5; page 154]
• The tendency of an atom covalently bonded in a molecule to attract
to itself the electron pair forming the covelent bond.
• The greatest electronegativity values are found with small
non!metals.....
Some Electronegativity Values
F
4.0
•
O
3.5
CR
3.2
Br,N
3.0
I
2.7
S
3.6
C
2.5
H
2.2
Na
0.9
Electronegativity is important in covalent bonding, because a bond
between two atoms of different electronegativity will be polarized
as a result.
EXAMPLES:
HCR:
– 130 –
H2O:
General Trend for Electronegativity in the Periodic Table
– 131 –
Electron Affinity
•
This is opposite of I.E; a gaseous atom gains an electron
M(g) +
eG
!
MG(g)
•
This process may be favourable (exothermic) or unfavourable
(endothermic).
Favourable values of E.A. are found when a valence shell is being
completed:
H(g) 1s1 + eG
!
H G (g)
1s2
-73 kJ
2
5
2
6
F(g) [He] 2s 2p
+ eG !
F G (g) [He] 2s 2p
-328 kJ
2
5
2
6
CR(g) [Ne] 3s 3p
+ eG !
CR G (g) [Ne] 3s 3p
-349 kJ
•
•
These anions are formed when the element meets a reactive metal,
e.g. Li+ HG.
•
Addition of a second electron is always unfavourable because of
like-charge repulsion:
O (g) + eG
! O G(g)
O G(g) + eG
! O 2G(g)
overall:
•
O(g) + 2eG
!
O 2G(g)
Nevertheless the O 2G ion is very common, as it is stabilized by
attraction to cations....
EXAMPLES:
In MgO: Mg2+ and O 2G
In BaO:
Ba2+ and O 2G
In H2O:
2 H + and O 2G
– 132 –